1 / 57

UNIT V

UNIT V. Electrolytic Cells. Electrolytic Cells. Electrolysis : uses an external power source to cause a non-spontaneous redox reaction to occur. In an ELC, the Anode is In an ELC, the Cathode is

yana
Télécharger la présentation

UNIT V

An Image/Link below is provided (as is) to download presentation Download Policy: Content on the Website is provided to you AS IS for your information and personal use and may not be sold / licensed / shared on other websites without getting consent from its author. Content is provided to you AS IS for your information and personal use only. Download presentation by click this link. While downloading, if for some reason you are not able to download a presentation, the publisher may have deleted the file from their server. During download, if you can't get a presentation, the file might be deleted by the publisher.

E N D

Presentation Transcript

  1. UNIT V Electrolytic Cells

  2. Electrolytic Cells Electrolysis : uses an external power source to cause a non-spontaneous redox reaction to occur. In an ELC, the Anode is In an ELC, the Cathode is connected to the POSITIVE connected to the NEGATIVE (-) (+) terminal of the external (-) terminal of the external power supply. power supply.

  3. Electrolytic Cells • In BOTH Electrochemical Cells (ECC’s) and Electrolytic Cells (ELC’s): OXIDATION takes place at the ANODE (LEOA) REDUCTION takes place at the CATHODE (GERC)

  4. Electrolytic Cells There are three main types of Electrolytic Cells: • Electrolysis of Molten Salts (no H2O) with Unreactive (Inert) Electrodes • Electrolysis of Aqueous Salts (H2O solution) with Unreactive (Inert) Electrodes • Electrolysis of Aqueous Salts (H2O solution) with Reactive Electrodes

  5. Type 1 - Electrolysis of Molten Salts (no H2O) with Unreactive (Inert) Electrodes • All molten (melted) salts consist of mobile ions. Eg.) NaI(l) Na+(l) + I-(l) NaCl(l) Na+(l) + Cl-(l)

  6. Type 1 - Electrolysis of Molten Salts (no H2O) with Unreactive (Inert) Electrodes • An example of a “Type 1” ELC: Electrolysis of molten NaCl (NaCl(l)) CATHODEANODE Reduction of Na+ Oxidation of Cl- happens here (GERC) happens here (LEOA)

  7. Type 1 - Electrolysis of Molten Salts (no H2O) with Unreactive (Inert) Electrodes To find the overall redox reaction with its Eo, we add up the half reactions as follows: (Na+ + e- Na(s)) 2 Eo = -2.71 v 2 Cl-  Cl2(g) + 2e-Eo = -1.36 v _______________________________________ 2 Na+ + 2 Cl-  2 Na(s) + Cl2(g)Eo = -4.07 v • The Product at the Cathode is Na(s) • The Product at the Anode is Cl2(g) • The industrial application of this cell is called a Down’s Cell. • In Electrolytic Cells, the Eo for the overall redox reaction is ALWAYS NEGATIVE. They are all non-spontaneous!

  8. Type 1 - Electrolysis of Molten Salts (no H2O) with Unreactive (Inert) Electrodes To Write Half-Reaction for Type 1 Electrolytic Cells: • Dissociate the salt into its ions. • Underneath the Cation, write C- (representing the Cathode). • Underneath the Anion, write A+ (representing the Anode). • Draw arrow showing where each ion goes. • The half-reactions must start with the ion (the ion must be the reactant). • For the Cathode Half-Reaction, write the reduction of the cation (same as on table). • For the Anode Half-Reaction, write the oxidation of the anion (reversed from table).

  9. Type 1 - Electrolysis of Molten Salts (no H2O) with Unreactive (Inert) Electrodes Example: a. Write the overall redox reaction for the Electrolysis of Molten Potassium Iodide (KI(l)):

  10. Type 1 - Electrolysis of Molten Salts (no H2O) with Unreactive (Inert) Electrodes b. Sketch this cell, labeling everything:

  11. Type 2 Electrolytic Cells – Electrolysis of Aqueous Salts & Non-reactive Electrodes • Because the salts are aqueous, water is present. Looking at the top and bottom shaded lines on the reduction table, you can see that water can be oxidized or reduced: The REVERSEof this reaction: H2O  ½ O2 + 2H+ + 2 e-Eo = -0.82 V is the oxidation of water and could take place at the ANODE. This half-reaction, the way it’s written: 2H2O + 2e-s H2 + 2OH- (10-7M) E° = -0.41 V is thereduction of water and could take place at the CATHODE.

  12. Type 2 Electrolytic Cells – Electrolysis of Aqueous Salts & Non-reactive Electrodes To Find the Half-Reaction at the CATHODE (Reduction): **Look at the “Overpotential Arrow” on the LEFT side of the table. 1. Is the Cation reduced ? OR 2. Is Water reduced?

  13. Type 2 Electrolytic Cells – Electrolysis of Aqueous Salts & Non-reactive Electrodes • All cations (+ ions) ABOVE the overpotential arrow WILL be reduced at the Cathode in aqueous solution (water present). Ex. Zn2+ + 2e- Zn(s) • All cations (+ ions) BELOW the overpotential arrow WILL NOT be reduced at the Cathode in aqueous solution. The Water will be reduced instead. 2 H2O + 2e- H2 + 2OH-

  14. Type 2 Electrolytic Cells – Electrolysis of Aqueous Salts & Non-reactive Electrodes • So when Cu2+ is in solution, what will be reduced at the cathode, Cu2+ or H2O? _____________ • Write an equation for the half-reaction at the Cathode: ___________________________________ Eo=_____So when Mg2+ is in solution, what will be reduced at the cathode, Mg2+ or H2O? ____________ • Write an equation for the half-reaction at the Cathode: ___________________________________ Eo=_____

  15. Type 2 Electrolytic Cells – Electrolysis of Aqueous Salts & Non-reactive Electrodes To Find the Half-Reaction at the ANODE (Oxidation): **Look at the “Overpotential Arrow” on the RIGHT side of the table. 1. Is the Anion oxidized ? OR 2. Is Water oxidized?

  16. Type 2 Electrolytic Cells – Electrolysis of Aqueous Salts & Non-reactive Electrodes • Any anion ABOVE the overpotential arrow WILL NOT be oxidized in aqueous solution. The water will be oxidized instead: 2 H2O  + ½ O2(g) + 2 H+ + 2 e- (reverse the reaction at +0.82 v) • Any anion BELOW the arrow WILL be oxidized in aqueous solution. : Ex. 2 Cl-  Cl2 + 2e- 2 Br-  Br2 + 2e-

  17. Type 2 Electrolytic Cells – Electrolysis of Aqueous Salts & Non-reactive Electrodes • So when SO42- is in solution, what will be oxidized at the anode, SO42- or H2O? _________ • Write an equation for the half-reaction at the Anode: _______________________________________ Eo=____ •  So when I- is in solution, what will be oxidized at the anode, I- or H2O? __________ • Write an equation for the half-reaction at the Anode: _______________________________________ Eo=____

  18. Type 2 Electrolytic Cells – Electrolysis of Aqueous Salts & Non-reactive Electrodes • To Write Half-Reaction for Type 2 Electrolytic Cells: • Dissociate the salt into its ions. • Write “H2O” between the two ions. • Underneath the Cation, write C- (representing the Cathode). • Underneath the Anion, write A+ (representing the Anode). • Using the rules about the “Overpotential Arrows” determine what is reduced at the Cathode (the cation or water) and what is oxidized at the Anode (the anion or water). • Complete the half-reactions at the Cathode and the Anode (with their Eo’s). • Add half-reactions to get the overall redox reaction.

  19. Type 2 Electrolytic Cells – Electrolysis of Aqueous Salts & Non-reactive Electrodes Example: A major industrial process is the electrolysis of brine (NaCl (aq)): inert anode inert cathode

  20. Type 2 Electrolytic Cells – Electrolysis of Aqueous Salts & Non-reactive Electrodes NaCl Na+ H2O Cl- C- A+ Now, we must look at the Reduction Table: Cathode Half-Reaction: 2 H2O + 2e- H2(g) + 2 OH-Eo = -0.41 V Anode Half-Reaction: 2 Cl-  Cl2(g) + 2e-Eo = -1.36 V Overall Redox Reaction: 2 Cl- + 2 H2O  Cl2(g) + H2(g) + 2 OH-Eo = -1.77 V

  21. Type 2 Electrolytic Cells – Electrolysis of Aqueous Salts & Non-reactive Electrodes • The products at the cathode would be H2(g) + 2 OH- (the pH near the cathode would ____crease). • The product at the anode would be Cl2(g). • The minimum voltage required to carry this reaction out would be 1.77 v (just enough to overcome the –1.77 v Eo).

  22. Type 2 Electrolytic Cells – Electrolysis of Aqueous Salts & Non-reactive Electrodes Example: • For the electrolysis of aqueous CuCl2 using platinum (inert) electrodes, find: • The half-reaction at the Cathode: _______________________________ Eo = ________ • The half-reaction at the Anode: _______________________________ Eo = ________ •  The overall redox reaction: _______________________________ Eo = ________

  23. Type 2 Electrolytic Cells – Electrolysis of Aqueous Salts & Non-reactive Electrodes • Product(s) at the Cathode: _________________ • Product(s) at the Anode: _________________ • The minimum voltage required: ________________ V

  24. Type 2 Electrolytic Cells – Electrolysis of Aqueous Salts & Non-reactive Electrodes Example: • For the electrolysis of Na2SO4(aq) using carbon (inert) electrodes, find: • The half-reaction at the Cathode: _______________________________ Eo = ________ • The half-reaction at the Anode: _______________________________ Eo = ________ •  The overall redox reaction: _______________________________ Eo = ________

  25. Type 2 Electrolytic Cells – Electrolysis of Aqueous Salts & Non-reactive Electrodes • Product(s) at the Cathode: _________________ • Product(s) at the Anode: _________________ • The minimum voltage required: _________________ V

  26. Type 2 Electrolytic Cells – Electrolysis of Aqueous Salts & Non-reactive Electrodes Example: • For the electrolysis of CuSO4(aq) using inert electrodes, find: • The half-reaction at the Cathode: _______________________________ Eo = ________ • The half-reaction at the Anode: _______________________________ Eo = ________ •  The overall redox reaction: _______________________________ Eo = ________

  27. Type 2 Electrolytic Cells – Electrolysis of Aqueous Salts & Non-reactive Electrodes • Product(s) at the Cathode: _________________ • Product(s) at the Anode: _________________ • The minimum voltage required: _________________ V

  28. Type 3 Electrolytic Cells – Electrolysis of Aqueous Salts with Reactive Electrodes • In this type of cell, the electrodes are normal metals, not inert ones like platinum or carbon. • The metal that the Cathode is made from is NOT reduced. That’s because metalscannot gain electrons and become negative metal ions (no such ion as Fe2-!) • The Cathode metal is NOT oxidized. That’s because oxidation does not take place at the cathode!

  29. Type 3 Electrolytic Cells – Electrolysis of Aqueous Salts with Reactive Electrodes THE CATHODE METAL NEVER REACTS!! • The cathode only supplies the SURFACE for the reduction of the cation (+ ion) or forthe reduction of water (depending on whether the cation is above or below the left overpotential arrow).

  30. Type 3 Electrolytic Cells – Electrolysis of Aqueous Salts with Reactive Electrodes Example: An aqueous solution containing the Cu2+ ion is electrolyzed. The cathode is made of iron. Write the equation for the half-reaction taking place at the cathode. Cathode Half-Reaction: Cu2+ + 2e- Cu(s) Eo = + 0.34 v

  31. Type 3 Electrolytic Cells – Electrolysis of Aqueous Salts with Reactive Electrodes Example: An aqueous solution containing the Mg2+ ion is electrolyzed. The cathode is made of nickel. Write the equation for the half-reaction taking place at the cathode: Cathode Half-Reaction: _________________________________ What would you observe at the cathode? _________________________________

  32. Type 3 Electrolytic Cells – Electrolysis of Aqueous Salts with Reactive Electrodes • In a Type 3 cell, there are THREE Possibilities for Oxidation at the ANODE: • The ANION in the solution is oxidized. • WATER is oxidized. • The ANODE METAL is oxidized. • The one with the highest oxidation potential (the LOWEST one on the RIGHT of the table) will be the one that is oxidized. Treat water as if it were at the right overpotential arrow.

  33. Type 3 Electrolytic Cells – Electrolysis of Aqueous Salts with Reactive Electrodes Example An aqueous solution containing the Cl- ion is electrolyzed. The anode is made of silver. Write the equation for the half-reaction taking place at the anode: Anode Half-Reaction: Ag(s)  Ag+ + e- Eo= - 0.80v

  34. Type 3 Electrolytic Cells – Electrolysis of Aqueous Salts with Reactive Electrodes Example:Given the following cell: Iron Nickel Al2(SO4)3

  35. Type 3 Electrolytic Cells – Electrolysis of Aqueous Salts with Reactive Electrodes Example: Identify: • The Anode: _______________ • The Cathode: _______________ • The half-reaction at the Cathode: _______________________________ Eo = ________ • The half-reaction at the Anode: _______________________________ Eo = ________ •  The overall redox reaction: _______________________________ Eo = ________

  36. Type 3 Electrolytic Cells – Electrolysis of Aqueous Salts with Reactive Electrodes • Product(s) at the Cathode: _________________ • Product(s) at the Anode: _________________ • The minimum voltage required: _______________ V

  37. Electrorefining • In electrorefining, an impure metal is refined by making it the ANODE. The pure metal is the CATHODE. IMPURE ANODE  PURE CATHODE (PC)

  38. Electrorefining Example: An ANODE made of IMPURE COPPER could have something like the following make-up: • Mostly copper • Some Ag or Au (above Cu on the reduction table) • Some Zn (below Cu on the reduction table) • Some “dirt” –any non-metal impurities The CATHODE would be made of PURE COPPER. The SOLUTION is an aqueous solution of a compound containing Cu2+ (eg CuSO4).

  39. Electrorefining • Refer to page 245 in Hebden Textbook

  40. Electrorefining SUMMARY: • The impure metal is the anode. • The pure metal is the cathode. • The electrolyte (solution) contains the cation of the metal to be purified. • Hebden Textbook page 246 Question #77

  41. Electroplating In Electroplating: • The object to be plated is the cathode (attached to the negative terminal of the battery) • The electrolyte must contain the cation of the metal to be plated on the object • The best anode is the metal to be plated onto the object

  42. Electroplating Example: We wish to plate a brass ring with copper. • The brass ring is made the cathode (connected to the negative) • The anode is copper (to keep supplying Cu2+ ions as it is oxidized) • The electrolyte is aqueous CuSO4 (supplies Cu2+ to start with - SO42- does not react here)

  43. Electroplating • Hebden Textbook page 244 Questions #73-76

  44. Electrowinning • The name given to the reduction of ores to produce metals in industry. • Ex) Zn2+ + 2e- Zn(s) Comes from zinc ores Zinc Metal like ZnS, ZnCl2 etc.

  45. Electrowinning Aluminum Production • Refer to page 246 of Hebden Textbook • Al3+CANNOT be reduced from an aqueous (water) solution • In order to reduce Al3+ to the metal Al, you must electrolyze a MOLTEN salt of Al • The main ore used to produce Al is called Bauxite – Hydrated Aluminum Oxide or Al2O3.3H2O

  46. Electrowinning • When Bauxite is heated, the water is forced off: Al2O3.3H2O + heat  Al2O3 + 3H2O “Alumina” • The melting point of Alumina (Al2O3) is much too high to melt it economically (mp = 2072 oC) • So Alumina (Al2O3) is mixed with Cryolite (Na3AlF6) • The melting point of this mixture is ≈ 1000 oC

  47. Electrowinning • Remember, there is NO water in this mixture. It is molten. • Some of the ions present in this mixture of molten salts are: Al3+, O2-, F- and Na+ • Cathode Half-Reaction: Al3+(l) + 3e- Al(l) • Anode Half-Reaction: C(s) + 2O2- CO2(g) + 4e-

  48. Electrowinning • This process is carried out by Alcan in Kitimat, B.C. • It uses 10 million amps of electricity. • Alcan has their own power generating plant at Kemano B.C.  • Because of high temperature and great amounts of electrical energy used, production of Al is expensive! Recycling Al uses a lot less energy!

  49. Electrowinning Alcan, Kitimat, B.C. • Hebden Textbook page 246 Questions #78-80

  50. Corrosion • Corrosion is undesirable oxidation of a metal (usually Fe).  • Oxygen from the air is reduced at the cathode. • The Fe becomes the anode (is oxidized). • Moisture can provide a solution for these processes to take place in. • O2 from the air dissolves in the waterdroplet. The dissolved O2 is reduced at the cathode and the iron surface becomes the anode.

More Related